The Atom Lesson 3 : The Bohr Model

Bohr Model of an Atom Electrons orbit the nucleus in fixed energy ranges called orbits (energy levels) An electron can move from one energy level to another by gaining or losing discrete amounts of energy Electrons cannot be found between energy levels (think of energy levels like rungs on a ladder . . . ) The lowest energy level is closest to the nucleus, the highest is farthest away The electron energy levels are quantized. Think of it like potential energy, the higher you are the more you have

Absorption Vs Emission When an electron (e-) absorbs (gains) energy (in whole photons or “quanta”) it “jumps” to a higher energy level This is called the EXCITED STATE When an e- emits (loses) energy it falls to a lower energy level and the energy emission is given off as photons (light) This is called the GROUND STATE The return to ground state is what we see as color in the flame test

So how was the “color” made in the flame test? Scientists use the Bohr model to explain this phenomenon Hydrogen Spectrum Flame Test

There is NO net change in energy Energy absorbed = energy released = energy of light produced Sometimes (like the flame test) this light is in the small section of wavelengths called the visible spectrum and we can see it Most of the time the human eye cannot

Bohr’s Hydrogen Model Turn to page 8 in your ref. packet When an electron falls from n=6 to n=3 what wavelength of light will be emitted? 1094 nanometers What region of the spectrum does that wavelength correspond to? Infrared Would we see it? Not with our naked eye

Hydrogen’s Line Spectrum Hydrogen emits four visible wavelengths of light Visible light is emitted when an excited electron “falls” from n= 3, 4, 5, or 6 back to n=2 434 nm; blue 656 nm; red Faint lines are UV 410 nm; violet 486 nm; cyan

Practice What color of light will be emitted if an e- goes from: n=6 to n=2? Violet (410nm) n=5 to n=2? Blue (434nm) n= 3 to n=2? Red/orange (656 nm)

Evidence for Energy Levels Bohr realized that this was the evidence he needed to prove his theory. The electric charge temporarily excites an electron to a higher orbit. When the electron drops back down, a photon is given off. The red line is the least energetic and corresponds to an electron dropping from energy level 3 to energy level 2.

Electromagnetic Spectrum (EM) EM is the complete range of electromagnetic radiation The spectrum includes most types of radiation, most of which are invisible to the human eye. The visible spectrum is the range of wavelengths between 400 and 700 nm. Chapter 5

Wavelength increases Frequency increases Energy increases

Anatomy of a Wave Wavelength ( λ ) – the shortest distance between equivalent points on a continuous wave. Wavelength is measured is units of length - m, mm, µm, nm Amplitude – the height from the origin to the crest (or trough) Frequency ( ν ) – the number of waves that pass a given point in one second

Wave Nature of Light Light travels through space as a wave, similar to an ocean wave. As frequency increases, energy increases (direct relationship) Wavelength (λ) and frequency (ν) have an inverse relationship…. As λ increases, ν decreases Short Wavelength = High Frequency = High Energy Long wavelength = Low Frequency = Low Energy

The longer the wavelength of light, the lower the frequency The longer the wavelength of light, the lower the frequency. The shorter the wavelength of light, the higher the frequency.

The Wave/Particle Nature of Light In 1900, Max Planck proposed that radiant energy is not continuous, but is emitted in small bundles. This is the quantum concept. Radiant/Light energy has both a wave nature and a particle nature. An individual unit of light energy is a photon.